Hydrogen | Group 1 - Alkali Metals | Group 2 - Alkaline Earth Metals | Group 12
6.941 22.9898 39.0983 85.467 132.9054 (223)
Sodium and potassium are the most abundant Group 1 elements on the Earth, and can be found in rock salt deposits formed from salt water evaporation. Lithium, rubidium and caesium are much less abundant. Francium is radioactive and only has short lived isotopes formed from the decay of other elements.
Sodium and potassium ions are very important in biology for the transmission of nerve impulses and maintaining osmotic processes. Several lithium salts are also used in psychiatric treatments. Compounds of sodium, including sodium hydroxide, sodium carbonate, sodium sulphates and sodium silicates, are very important industrial chemicals and sulphates of potassium are widely used in fertilisers.
Preparation of the Elements and Their Physical Properties
Lithium and sodium metal can be prepared by the electrolysis of molten salts of these elements. This is not as effective for potassium, rubidium and caesium, so their isolation is carried out by treating their molten chloride salts with sodium vapour. The metals are silvery in appearance, apart from caesium which is golden yellow. Because the metals all have only one valence electron, lattice bonding is fairly weak and so the metals are soft - they can be easily cut with a knife - and have low melting points. They will react with many elements if heated.
The chemistry in group 1 is dominated by the metal +1 ions. +2 ions of the elements are not easily formed due to their very high second ionisation energies. The compounds are essentially ionic in nature, although there is some degree of covalent character for some Li+ compounds. As the size of the M+ ions get larger and their charge density reduces, the chemistry becomes more ionic in nature.
The metals all react with water to produce hydrogen gas and the metal hydroxide. Lithium only reacts very slowly; sodium much faster; potassium with burning; rubidium and caesium explosively. Sodium, and sometimes potassium, are used therefore used to remove traces of water from hydrocarbon solvents, such as hexane or benzene, for use in air and moisture sensitive reactions.
M + H2O → MOH + ½ H2
|Melting point (°C)
|Ionic Radius (Å)1
|Ionisation enthalpy (kJ mol-1)
|Approx hydrated radius of M+ (Å)
|Approx hydration number M+
They also react with oxygen to form oxides. Li, Na and K are usually stored in mineral oil to avoid their surfaces from becoming tarnished. Rubidium and caesium require storage under an inert argon atmosphere. Different group 1 metals form different types of oxide. Lithium forms an oxide with the oxide ion O2- to form Li2O. Sodium forms an oxide with the peroxide ion O22- to form Na2O2. When heated under pressure, however, this will form Na2. The other alkali metals - K, Rb and Cs - form oxides with a superoxide ion O2- to give MO2.
Na + CH3CH2OH → CH3CH2O-Na+ + ½ H2
Another common use of sodium in the laboratory is in sodium amalgam. The metals are all soluble in mercury and Na/Hg is a powerful reducing agent.
One of the more interesting properties of the group 1 elements (and also of calcium, strontium, barium, europium and ytterbium) is that they will dissolve in ammonia to give solutions, which, when dilute, are deep blue in colour. What happens is that, when the metals dissolve, they are ionised, releasing the single valence electron. The electron is solvated and interacts with ammonia molecules, resulting in the deep blue colour. The metal ions M+ and the metal atoms are colourless.
Na → Na+ + e-
As these solutions become more concentrated, this blue colour gives way to a more copper coloured, almost metallic looking, liquid. At this stage, clusters of metal ions form and these solution have very high electrical conductivity.
The blue solutions can be quite stable over time but addition of certain transition metal compounds or use of bright light can give the formation of sodium amide, NaNH2 which can be used as a powerful reducing agents and strong bases.
Na + NH3 → NaNH2 + ½ H2
Solvation of the Metal Ions
When an ionic compound of an alkali metal is dissolved in a solvent, it can't be considered to be a bare ion. The metal ion has a number of solvent molecules associated with it that surround the ion and whose behaviour is affected by it. This is called the solvent shell, and if the solvent is water, the hydration shell.
It can be divided up into two areas:
The primary solvent shell - the group of solvent molecules that are in direct contact and are coordinated to the metal.
A secondary solvation shell - molecules that are outside the primary shell but whose behaviour and arrangement are still under the influence of the metal ion.
The total number of water molecules under the influence of an ion in aqueous solution is called its hydration number (table 1). For Li+, Na+ and K+, the primary hydration shell consists of four water molecules, and, for Rb+ and Cs+, there are probably six. Although the size of the primary hydration shell increases going down the group due to the greater size of the ions, the overall hydration number reduces and so does its hydrated radius - the size including the whole solvent shell.
This is because, as the size of the ion increases, its charge density reduces, giving it a smaller influence on the solvent around. This difference in size of the hydrated ions, as well as some other factors, is of vital importance to the ability of our body's cells to discriminate between different ions.
Alkali Metal Compounds
The formation of the oxides of these elements by reaction with excess O2 has already been mentioned above. The hydroxides, MOH, are all white crystalline solids, readily soluble in water, and are hygroscopic (will absorb water vapour from the air and appear visibly wet). The solids and solutions of the hydroxide all absorb carbon dioxide and LiOH is used as a scrubbing agent to remove deadly CO2 from the air inside space craft.
The elements form ionic salts of virtually every acid and are usually colourless crystalline solids. Ionic slats of these elements generally have higher melting point, and, once melted, are electrically conductive and water-soluble. Many lithium compounds have properties that differ from the rest of the alkali metals; more like those of Mg2+ compounds. For example, LiH is stable up to about 900 °C, while NaH decomposes at a much lower temperature of about 350 °C.
The compound Li3N, formed from the reaction of lithium with nitrogen gas at high temperature, is stable. The corresponding sodium analogue, Na3N, however, does not exist at room temperature. LiOH is much less soluble than any of the other hydroxides and also decomposes to Li2O at high temperature whereas the others sublime. The lithium salts of strong acids tend to be the most soluble of the alkali metals in water, whereas the salts of weak acids tend to be the least soluble.
Complexes of Alkali Metal Ions
As described, the alkali metal ions will have a significant interaction with the solvent it is in. Alkali metal ions can also form complexes, though not quite in the same way as transition metals.
Compounds called 'crown ethers' and 'cryptands' can encapsulate these metal ions and will allow them to enter solvents in which they would otherwise be insoluble. Crown ethers are organic molecules and are classed as 'cyclic polyethers'. These ring molecules have oxygen atoms that can coordinate with the ion spaced by short stretches of CH2 units and have the general form (CH2CH2O)n, where n is the number of these units that make up the ring.
One of the most common crown ethers is called 18-crown-6 (see figure 1). The '18' denotes that there are 18 atoms making up the backbone of the ring; the word 'crown' reveals that it is a crown ether and the number '6' tells us the ring contains six oxygen atoms: n is 6.
These molecules can be selective in binding to different alkali metal ions. Different sized ions are selectively bound by different sized rings. For instance, 18-crown-6 binds Li+ < Na+, Cs+ < Rb+ < K+ so in a mixture of the various ions, it will selectively bind potassium.
Figure 1: Structure of the crown ether 18-crown-6 H2C-CH2 / \ H2C-O O-CH2 / \ H2C CH2 | | O O | | H2C CH2 \ / H2C-O O-CH2 \ / H2C-CH2
Cryptands are similar in their complexation of alkali metal ions but are more selective. They also differ in that they contain nitrogen atoms to coordinate to the metal as well as oxygen. Cryptands contain two nitrogen atoms that are connected by three ether chains. A common cryptand is 2,2,2-crypt and has the structure N(CH2CH2OCH2 CH2OCH2CH2)3N. This encapsulates the metal ion coordinating to it through all six oxygen atoms and both nitrogens.
The use of this ligand led to the discovery of some very bizarre compounds. When 2,2,2-crpyt was added to a cooled solution of sodium in ethylamine, a compound crystallised with the formula [Na(2,2,2-crypt)]+Na-. It is only stable below -10 °C. Such compounds containing negatively charged sodium ions are called 'sodides'. There are also examples of compounds, known as 'electrides', where the sodide ion is replaced by an electron - for example [Cs(crypt)]+e-.
These types of encapsulating ligand imitate the behaviour of cyclic polypeptides that play a role in the transport of alkali and alkaline earth metal ions across cell membranes.
Organometallic Compounds of Alkali Metals
Organolithium compounds have great importance to organic chemistry as sources of R-, where R is an organic group (such as an alkyl group). RLi compounds are soluble in most organic solvents, can be liquids and are very reactive and sensitive to moisture. Most are spontaneously flammable in contact with air and need careful handling. These are best prepared by the direct reaction between lithium metal and an organic halide, for example:
2 Li + C2H5Cl → C2H5Li + LiCl
Most organolithium compounds have significant covalent character. Methyllithium, CH3Li, in the solid state consists of tetrahedral arrangements of lithium atoms with four methyl groups occupying the faces of this tetrahedron. The methyl groups form four centre-two electron covalent bonds with the three Li atoms of the face it occupies.
Sodium and potassium also form organometallic compounds, but these are essentially ionic and are sparingly soluble in organic solvent. They are, however, exceedingly reactive. The metal react can react with organic molecules in which certain hydrogen atoms are slightly acidic. For example, sodium react with cyclopentadiene and terminal alkynes
3 C5H6 + 2 Na → 2 C5H5-Na+ + C5H8
RC≡CH + Na → RC≡C-Na+ + 1/2 H2
Both are very important reagents and ligands in the synthesis of transition metal organometallic complexes, especially C5H5-Na+.